Honor S Chemistry: Fall Semester Final

Honor’s Chemistry: Fall Semester Final Name______

1. Explain what is wrong with the statement “My friend burned a piece of paper (a hydrocarbon) that had the final exam on it and it disappeared”. (Be sure to use a chemical equation, identify reactants and product(s) and include energy).

The paper (CxHy) was burned with oxygen and the atoms in the paper were broken apart and rearranged into new combinations. The new combinations were the products CO2 + H2O Carbon dioxide and water are always the products of a combustion reaction of a hydrocarbon. The paper didn’t disappear, but its atoms were rearranged into gases.

Law of conservation of mass – mass is not created nor destroyed...atoms are only rearranged in chemical reactions. [Reactants = Products]

CxHy + O2 --> CO2 + H2O + energy

REACTANTS PRODUCTS

Reaction is EXOTHERMIC

match was activation energy (EA) to start reaction R Energy P

2. Write a balanced chemical equation for the reaction forming magnesium chloride precipitate from its elements. Draw a picture to help me visualize what is happening.

Mg (s) + Cl2 (g) --> MgCl2 (s)

Mg C C C Mg C l l l l

3. Identify four specific errors made during experiments that would disobey the scientific method.

 have bias in conclusion  do not use a control for comparison  make measurements a single time or multiple times (with poor precision)  change two or more variables at a time  exclude data because it doesn’t fit with the rest  not make careful observations  make-up data SCIENTIFIC METHOD  Observation  Hypothesis  Collect Data  Analysis  Conclusions  Repeat / Modify

4. Explain which is more meaningful data: quantitative or qualitative.

5. Compare and contrast the terms theory and law.

Natural law – describes events in nature laws do not change laws of nature will always occur and are not man-made

Theory – an explanation of an event theories can change as new evidence is discovered. theories are man-made

6. Explain how the scientific method is used to solve a problem. Describe the design of a controlled experiment. 7. Describe to a General Chemistry student how to make a measurement correctly.

8. Report in scientific notation, using 3 significant figures, the number 670,842,004.

9. Convert the following: a. 100 mm  x dm

b. 3 x 108 m/s  x km/day

c. 2 x 10-8 cm  x m

10.Your friend tells you that the number 1.2000 is more accurate than 1.2 x 100. Is your friend correct? Explain.

11. If 200 cops = 4 criminals and 3 cows = 1 horse and 4 horses = 9 cops. How many cows equal 1000 criminals. Show work.

12. Your friend says that smoking a mercury-laced cigarette is cool. You aren’t convinced and decide to look up the LD50 value of mercury. It is 0.4 mg/kg. Assuming you weigh 150 lbs and that 2.2 lb = 1 kg. How much mercury can you safely smoke?

13. List 4 intensive properties and 3 extensive properties of a BabyRuth candy bar. INTENSIVE EXTENSIVE

14. Draw and label a phase diagram for a non-pure substance that has a melting point of 22oC and a boiling point of 89oF. 15. Which has more kinetic energy a 400 mg bullet moving at 250 m/s or a lead ball, moving at 0.01 m/s. The radius of the lead ball is 30 dm and the density of lead is 11.2 g/cm3. [V = 4/3  r3]

16. Justify why or why not we should pursue an energy program of nuclear fusion in the United States. You need to explain the differences in fission and fusion, site advantages and disadvantages of each. 17. Classify the following materials as elements, compounds, or mixtures: a. Lead (II) chloride b. ozone c. vinegar d. heavy water e. tin foil

18. How much heat will be absorbed by a 20 g piece of ice (at 254 K) that is warmed to 150oF? Latent heat of vaporization (H2O) = 2256 J/g Latent heat of fusion (H2O) = 333 J/g Specific heat of water (liquid) = 4.184 J/goC Specific heat of water (solid) = 2.077 J/goC Specific heat of water (gas) = 2.042 J/goC

19. What is the final temperature of a 20 g block of ice (at 254 K) that is placed in 300 g of water (T = 50oC) 20. Explain how Archimedes principle would be used to determine if a gold crown was “pure” gold. What other information would you need to know to be certain? 21. Summarize the development of the current atomic theory.

22. Given the following: U-235

a. Write the longhand and shorthand electron configuration for U-235

b. How many protons ___, neutrons___, and electrons___ does the element have?

c. Write the formula for an isotope of this element. 23. Given that light has a wavelength of 412 nm. What is its energy?

24. Draw an energy level diagram for an Al3- anion. Be sure to explain how the Aufbau principle, Pauli exclusion principle and Hund’s rule have been obeyed.

25. Draw a diagram to explain the lyman, paschen and balmer series in the hydrogen atom. Include a brief explanation. 26. Compare and contrast the terms ions, atoms, and isotopes in subatomic structure.

27. Calculate % a mass of an isotope “X”: given that the average atomic mass of “X” is 54.3 g/mol and the element has only two isotopes (X-50 comprises 38% abundance).

28. Why do metals have lower ionizations energies than nonmetals?

29. What differences in atomic structure (microscopic) explain the observable (macroscopic) differences in salts and compounds made from two non-metals? 30. What is for formula for the following: a. Ammonium chloride

b. Diphosphorous octabromide

c. Plumbous cyanide

31. a) Draw the Lewis structure for the phosphite ion.

b) What is the apparent charge on the P atom?

c) What is the percentage composition in ammonium nitrite?

32. Empirical molecular formula 33. mole island

34. polar vs. non-polar

35. Which has more atoms: 396 g titanium (II) sulfate, 2.3 x 1022 molecules trichloro 3 nonaoxide or 4.5 x 10 dm3 of methane (CH4) gas @ STP? Show work for credit.

Balance the following equation:

Mole  mol G  liters Molecules  mole G  atoms 15. Describe the difference between a natural law and a theory.

36. Suppose that you attempt to turn on a lamp, but the bulb does not light. Using the scientific method, describe how you might solve this problem. Be as complete as you can, and identify the elements of the scientific method in your explanation. 37. The substance looked pale yellow and had a density of 3.6 g/mL. It burned readily in air, and produced bubbles when reacted with acid. When heated, it changed from solid to liquid at 79oC, and from liquid to gas at 143oC.

38. Identify the following properties as either chemical or physical

a. pale yellow

b. density of 3.6 g/mL

c. burned readily in air

d. produced bubbles when reacted with acid

39. Is there any difference between the properties of pure water that has been boiled and condensed and the properties of pure water that has been frozen and then melted? Explain

40. You are given a flask that contains sea water that has been contaminated with oil. Some sand is also present in the flask. Describe how you would separate the sand, oil, sea salt, and water from each other.

41. Complete the following table: Element Symbol Atomic No. of No. of Mass No. of Charge (atom/ion Number protons Neutrons Number electrons ) 17 18 18

1 0 +1

23 sodium Na atom 11

42. Write the formula for the compounds that would be formed from the following ions:

Na+ and Cl- ______

Al3+ and Br- ______

K+ and S2- ______

Mg2+ and Cl- ______

43,. Compare Rutherford’s model of the atom to Thomson’s model. Explain Rutherford’s reasoning in developing his model

44.How might the results of Rutherford’s experiment have been different if he had used aluminum foil (atomic number 13) rather than gold foil (atomic number 79)? 45. a. aluminum sulfide ______b. SF2 ______c. phosphorus trichloride ______d. Zn(NO3)2 ______e. iron(III) oxide ______f. CuI ______g. HNO3 ______h. aluminum hydroxide ______i. CaBr2 ______j. hydrochloric acid ______k. Ba3(PO4)2 ______l. magnesium sulfite ______m. LiC2H3O2 ______n. nitrogen trichloride ______o. CuSO3 ______p. sodium carbonate ______

46. Round each number to the indicated number of significant figures and express it in scientific notation: a. 2501 (2 S.F) ______b. 0.030490 (3 S.F) ______c. 172590 (1 S.F.) ______d. 40035.2 (2 S.F) ______

47. a 12.6 m x 2.0 m x 13.84 m = ______b. 13 cm + 10.4 cm + 1.25 cm = ______c. (1.360 x 105 cm) x (6.05 x 10-2 cm) = ______d. 11.63 mL – 8.8 mL = ______e.. (12.36 g – 11.25 g) = ______10.31 mL f. 18.5 m = ______0.035 s

48. If 4 quarts = 1 gallon, and 1.06 quarts = 1 liter, how many liters are there in a 55.0 gallon container?

49.To three significant figures how many seconds are there in exactly 1 “microyear”?

50.On Nov. 20, 1977, Kenneth Peter Warby attained a speed of 345.48 miles per hour in his unlimited hydroplane Spirit of Australia. How fast was this in meters per second? (1 mile = 1.6093 km)

51. Calculate the mass percent of each element in (NH4)2S.

24 52. Find the mass, in grams, of 2.65 x 10 molecules of Cl2. 53. How many grams of sulfur are present in 83.2 g of sulfur dioxide?

54. How many hydrogen atoms are in 52.0 g of water?

55. Determine the empirical formula for a compound that contains 14.7 g of nickel and 40.0 g of bromine.

56. Balance the following chemical equations:

a. PbI2 + AgNO3 à Pb(NO3)2 + AgI

b. Mg + TiCl4 à MgCl2 + Ti

c. C3H8 + O2 à CO2 + H2O

d. P4 + O2 à P4O10 e. Na2CO3 + HCl à NaCl + CO2 + H2O

57. Write balanced equations for the following reactions: a. zinc + hydrochloric acid à zinc chloride + hydrogen (gas)

b. barium chloride + ammonium sulfate à barium sulfate + ammonium chloride

c. calcium hydroxide + nitric acid à calcium nitrate + water

d. calcium carbonate + hydrochloric acid à calcium chloride + carbon dioxide + water

e. bromine + sodium iodide à sodium bromide + iodine

f. magnesium + iron(III) chloride à magnesium chloride + iron

59. Lithium metal reacts with water to form aqueous lithium hydroxide and hydrogen gas.

60. Write a balanced chemical equation for the reaction, including abbreviations for the physical states.

60. Iron(III) nitrate in water solution reacts with potassium sulfide in water solution to form aqueous potassium nitrate and solid iron(III) sulfide. Write a balanced chemical equation for the reaction, including abbreviations for the physical states. 61. Potassium chlorate (KClO3) decomposes to form potassium chloride and oxygen gas. If 5.4 moles of potassium chlorate decompose, how many moles of oxygen could be produced?

62.What mass of FeCl2 could be produced from 35.0 g of Fe and excess HCl if the balanced reaction is

Fe + 2 HCl à FeCl2 + H2

63. When ammonia burns in pure oxygen, the reaction is:

4 NH3 + 3 O2 à 2 N2 + 6 H2O

64. What masses of nitrogen and water could be produced from 45.0 g of ammonia?

65. Copper metal reacts with a solution of silver nitrate, AgNO3, to produce copper (II) nitrate and silver metal. In carrying out this reaction, a piece of copper wire was immersed in a solution of silver nitrate until the reaction stopped. The original mass of the copper wire was 2.36 grams. After the reaction stopped, the mass of the wire was 1.03 grams. What mass of silver was produced?

66. If a piece of aluminum of mass 4.50 g and temperature 99.5oC is dropped into 12.0 g of water at 21.0oC, what will be the final temperature of the water-aluminum mixture? The specific heat capacity of aluminum is 0.902 J/(g. oC).

67. Write electron configurations for each of the following. DO NOT use noble gas shorthand.

68. Using noble gas shorthand, write the electron configuration for astatine (At). 69. Identify the elements that have the following electron configurations. If the configuration shows the atom in an excited state, write the ground state configuration for the atom.

a) 1s22s22p63s23p2

b) 1s22s22p63s23p64s23d104p4

c) 1s22s22p33p1

70. Using atomic structure in your explanation, account for the general trend in atomic size as you go from left to right across a period and from top to bottom down a group on the periodic table.